Lewis Concept of Acids and Bases

 

Gilbert Newton Lewis (1875-1946)

Gilbert N. Lewis was an American physical chemist who made significant contributions to the understanding of chemical bonding and reactions. He was born on October 23, 1875, in Weymouth, Massachusetts, and he passed away on March 23, 1946.

Lewis is best known for developing the Lewis dot structures, which are diagrams that show the bonding between atoms in a molecule and the lone pairs of electrons that may exist. These structures help chemists visualize and understand how atoms are connected in molecules.

In addition to the Lewis dot structures, he proposed the Lewis acid-base theory. According to this theory, an acid is a substance that can accept an electron pair, and a base is a substance that can donate an electron pair. This expanded the understanding of acids and bases beyond the traditional definitions.

Gilbert N. Lewis had a successful academic career, teaching at Harvard University and later at the University of California, Berkeley. He received numerous honors for his work, including the Langmuir Award in 1934.

His work laid the groundwork for the modern understanding of chemical bonding, and his concepts, such as Lewis dot structures and the Lewis acid-base theory, continue to be fundamental in the field of chemistry.

 This theory is based on electron pair acceptance and donation. In this concept, any compounds capable of accepting lone pairs of electron are considered as acids while the compounds that are capable of donating electron pair are described as bases. In this case when an acid and base reacts with each other a co-ordinate bond forms where a base donates its electrons to an acid.

NH₃+BF₃ → H₃N-BF₃

The relation between Lewis concept with Bronsted-Lowry and Arrhenius concept can be seen very easily. Suppose an acid and base in Lewis acid-base concept is defined as substance that can accept and donate electron pair. Therefore an ionic complex build with both cation and anion can be deemed as an acid and base respectively. For Bronsted and Arrhenius concept an acid is represented as HA which in aqueous solution generate H⁺ and A^- ions. Now Bronsted specifically mentions bases are capable of accepting protons for which it needs lone pair of electron to bind. Therefore a Lewis base can bind with not only with a proton but with other cations also. So we can say that a Bronsted base is a Lewis base but a Lewis base is not just confined as Lewis bases. We can also say that Arrhenius and Bronsted acid that can be represented by HA, is actually an adduct of Lewis acid H⁺ and Lewis base A^-. There are several compounds that can act as a Lewis base by donating π electrons like C₂H₄, C₂H₂ etc, but cannot act as a Bronsted base.

All kind of cations which have a tendency to accept electron pairs are Lewis acid.

Ø  Smaller cations with higher positive charge density like H⁺, Li⁺, Be²⁺ etc are likely to act as stronger Lewis acid due to their tendency to accept electrons.

Ø  Larger cations of alkali and alkaline earth metals behave as relatively weaker Lewis acid due to their lesser tendency to attract electrons.

Ø  Higher charged cations have higher tendency to accept electrons relative to the lower charged cation of a given metals.

Ø  Therefore Fe³⁺ acts as a stronger Lewis acid than Fe²⁺. Sn⁴⁺ stronger than Sn²⁺ etc.

Electron deficient compounds where the central atom or molecules have accessible vacant electron orbitals can also acts like a Lewis acid.

Ø  BF₃, AlCl₃, GaBr₃ have B, Al, Ga as central atom which in general has electron deficiency. These elements usually form complexes that do not fulfill their octet like the above mentioned complexes. Therefore these complexes tend to accept anions or electron from electron rich compounds to complete their octet.

Ligands as Lewis bases:

When metals and ligands interact to form complexes, the ligands play the role of Lewis bases, and the metal centres act as Lewis acids. There are different types of Lewis base ligands:

1. σ-Base Ligands: These ligands can only donate the lone pair through σ-bonding. Examples include NH₃.
2. π-Donor or π-Base Ligands: These ligands can donate the electron pair through both σ- and π-bonds. Examples are H₂O, OH^-, and halides.
3. π-Acceptor or π-Acid Ligands: These ligands can donate the electron pair through a σ-bond but can also receive electrons back from the metal center through π back bonding. Examples include CO, CN^-, C₂H₄, and PR₃. These ligands form stable complexes with lower valent metal ions, such as Ni(CO)₄ and Cu(CN)₄^3-.

Certain molecules like H₂, N₂, and O₂ can also act as π-acid ligands.

Strength of Lewis acids and Bases:

Here it is important to mention that in terms of Lewis concept, the Lewis acid/base strength of a species is not its inherent property, but it largely depends on the nature of counterpart with which the adduct formation occurs. The stability of an adduct is controlled by many factors and the matching of hard-soft character is one of the contributing factor. The following Lewis acid strength runs for the hard Lewis bases (e.g. O-donor).

BCl₃ > AlCl₃ > GaCl₃

But, for the soft Lewis bases like dimethylsulphoxide (S-donor), the reverse order is true and the Lewis acid strength sequence runs as:

GaCl₃ > AlCl₃ > BCl₃

Lewis acid-base interaction through oxidative addition: In some cases, during the Lewis acid-base adduct formation, oxidation may take place (i.e. oxidative addition). These are:

R₃P: + O ⇋ R₃P: →O, O + :SO₃^2- ⇋ O←SO₃^2- (i.e. SO₄^2-)

S + :SO₃^2- ⇋ S←:SO₃^2- (i.e. S₂O₃^2-)

Expansion of valence leading to ate- and onium-complex: when a Lewis acid expands its valency to combine a negative ion, it leads to ate complexes.

Me₃B +LiMe → Me₄B^-Li⁺ (ate complex)

Similarly, onium salts are produced when a Lewis base expands its valency.    

     Me₃N + MeI → Me₄N⁺I^- (onium salt)

Merits of Lewis Acid-Base concept

Lewis acid-base theory, proposed by Gilbert N. Lewis, provides a broader and more versatile understanding of acid-base interactions compared to other traditional theories like the Arrhenius and Brønsted-Lowry theories. Here are some merits or advantages of the Lewis acid-base theory:

General Applicability:

The Lewis theory is more general than the Brønsted-Lowry and Arrhenius theories. It extends the concept of acidity and basicity beyond proton transfer in aqueous solutions to include a wide range of reactions in various solvents.

Inclusion of Non-Aqueous Systems:

Unlike the Arrhenius and Brønsted-Lowry theories, which are primarily focused on aqueous solutions, the Lewis theory is applicable to non-aqueous solvents and reactions involving substances other than protons (H⁺).

Coordination Chemistry:

Lewis acid-base theory is particularly valuable in coordination chemistry. It helps explain the formation of coordination complexes, where metal cations act as Lewis acids by accepting electron pairs from Lewis bases (ligands).

Versatility in Reaction Types:

Lewis acids and bases can participate in a wide variety of reactions, including not only acid-base reactions but also complexation reactions, redox reactions, and more. This versatility makes the Lewis theory more adaptable to different chemical scenarios.

Molecular Orbital Considerations:

Lewis acid-base interactions can be rationalized in terms of molecular orbitals. This allows for a more detailed understanding of the electronic structure changes during chemical reactions.

Predictive Power:

The Lewis theory provides a basis for predicting and explaining the reactivity of various substances in different chemical reactions. It helps chemists understand and predict the behavior of molecules in terms of electron-pair interactions.

Global Perspective:

Lewis acid-base theory provides a global perspective on chemical reactions, emphasizing electron flow and electron pair interactions. It allows chemists to view reactions more holistically, considering the movement of electron pairs rather than just protons.

Compatibility with Quantum Chemistry:

    Lewis acid-base concept is compatible with modern quantum chemistry, allowing for a more sophisticated and accurate theoretical description of chemical phenomena.

Limitations of Lewis acid

Ø  The Lewis acid-base theory, although helpful, has some drawbacks:

Ignores the Effect of Solvents: It doesn't take into account how the presence of different liquids can affect the behavior of acids and bases.

Not Comprehensive for All Acids and Bases: The theory mainly looks at how electrons are shared, missing out on some types of acid-base reactions that involve passing around protons or charged particles.

Can't Predict Strengths Precisely: It doesn't offer a precise way to predict how strong an acid or base is. While it talks about electron pairs, it doesn't give a clear measure of how acidic or basic something is.

Focuses Too Much on Electron Pairs: It strongly emphasizes the exchange of electron pairs, possibly overlooking other important factors in acid-base reactions.

Limited to Certain Compounds: It might not work for all types of substances. Even though it explains the behavior of many things, there are cases where other theories make more sense.

Doesn't Consider Size and Shape: It doesn't think about how the size and shape of molecules might affect how they react. In some situations, this can be important.

Doesn't Cover Redox Reactions: It doesn't talk about reactions where electrons are transferred. Sometimes, these reactions happen alongside the Lewis acid-base reactions.

In short, while the Lewis acid-base theory is good for understanding many things, it's not perfect. We need to use it carefully and know that in some situations, other theories might work better.

Reference

1) Concise inorganic chemistry by J. D. Lee.

2) Inorganic Chemistry by James E. Huheey, Ellen A Keither, Richard L. Keither, Okhil K. Medhi.